2.39      describe simple tests for the gases:

Gas	Test	Results
i    hydrogen (H2)	introduce lighted splint	explodes with squeaky pop
ii   oxygen (O2)	introduce glowing splint	reignites the splint
iii  carbon dioxide (CO2)	bubble through lime water	goes milky white*
iv   ammonia (NH3)	introduce moist red litmus paper	turns blue
v    chlorine (Cl2)	introduce moist indicator paper (any colour)	indicator paper is bleached white

Note:    The limewater goes milky because a white precipitate of insoluble calcium carbonate is formed. The equation for limewater (calcium hydroxide) reacting with carbon dioxide is:

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

2.31    deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

A reactive metal will displace a less reactive metal from its oxide.  These reactions are often known as thermit reactions.  A good example is the reaction between aluminium and iron (III) oxide:

2Al(s)  +  Fe2O3(s)  ®   Al2O3(s)    +   2Fe(s)

These reactions are very vigorous and exothermic once they get started (a special fuse is needed to start them).  In the example above, the temperature rises above 2000^OC and the iron and aluminium oxide are formed in the molten liquid state!  The further apart in reactivity the metals are, the more violent the reactions become, so Mg and CuO would be very dangerous!

Similarly, a more reactive metal will displace a less reactive metal from one of its salts.

e.g. 1    zinc will displace copper from a solution of copper (II) sulfate

            Zn(s) + CuSO4(aq)   →   ZnSO4(aq) + Cu(s)

In the above reaction we see the solution go from blue to colourless and the grey zinc gets covered in brown copper. These displacement reactions are exothermic.

e.g. 2    iron cannot displace magnesium from a solution of magnesium sulfate, since Fe is less reactive than Mg.

2.30     describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and copper

Metal	Air	Water	Dilute acid (HCl)
potassium	burns easily	violent with cold water	violent
sodium	burns easily	very fast with cold water	violent
lithium	burns easily	fast with cold water	very fast
calcium	burns easily	quite fast with cold water	very fast
magnesium	burns easily	steam needed	fast
zinc	slow, heat needed	steam needed	reacts steadily
iron	slow, heat needed	steam needed	reacts slowly
copper	slow, heat needed	no reaction	no reaction

Reactivity series

2.29     understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

These metals are listed in decreasing reactivity.  The more reactive the metal, the more quickly it will react.  Learn this reactivity series.  Can you make up some sort of funny sentence with names or symbols to help you to remember it?

2.28     describe a physical test to show whether water is pure.

Pure water boils at 100 ^OC and freezes at 0^OC.  You can check these to confirm the identity of the liquid.

2.27     describe the use of anhydrous copper(II) sulfate in the chemical test for water

Anhydrous copper(II) sulfate is white. It turns blue, forming the hydrated salt, when water is added.  This is an exothermic reaction.

CuSO4(s) + 5H2O(l) → CuSO4.5H2O(s)

2.26     describe the combustion of hydrogen

Hydrogen burns explosively with an orange flame, forming water.

2H2(g) + O2(g) → 2H2O(l)

This is the basis of the “squeaky pop” test for hydrogen gas.

Hydrogen and water

2.25     describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

All react to give hydrogen and a solution of the metal salt.  Magnesium is the most reactive, followed by aluminium, zinc, and iron.  We see bubbles of hydrogen gas formed: quickly for Mg and getting slower for the other metals.  All the solutions of the salts formed are colourless, apart from iron.  Iron (II) salts in solution are pale green.  (Note that aluminium has a protective layer of its oxide on the surface and will only react quickly if this is removed first.)

·        With hydrochloric acid, solutions of the metal chlorides are formed:

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g)            (remember that Al forms a ion with 3+ charge)

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

Fe(s) + 2HCl(aq) → FeCl2(aq) + 3H2(g)              (notice that iron (II) chloride is formed here)

·        With sulfuric acid, solutions of the metal sulfates are formed:

Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g)

2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g)

Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)           (notice that iron (II) sulfate is formed here)

2.24   understand that carbon dioxide is a greenhouse gas and may contribute to climate change.

The greenhouse effect is where molecules of compounds in the atmosphere absorb some of the heat radiated from the Earth’s surface and stop it escaping into back space.  Carbon dioxide is one of these compounds.  Increased amounts of CO₂ in the atmosphere could trap increasing amounts of heat, contributing to global warming and changing the climate.

2.23     explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

Carbon dioxide is dissolved in drinks at pressures higher than atmospheric pressure to make them fizzy.  Gases become more soluble in water as the pressure increases.

Since carbon dioxide is denser than air (and does not support combustion), it will settle on fires at ground level and exclude the less dense oxygen, extinguishing the fire.